ake a sample from the mixture for
analysis.]

[Note 2: When aqueous solutions of ferrous compounds are heated in the
air, oxidation of the Fe^{++} ions to Fe^{+++} ions readily occurs in
the absence of free acid. The H^{+} and OH^{-} ions from water are
involved in the oxidation process and the result is, in effect, the
formation of some ferric hydroxide which tends to separate. Moreover,
at the boiling temperature, the ferric sulphate produced by the
oxidation hydrolyzes in part with the formation of a basic ferric
sulphate, which also tends to separate from solution. The addition of
the hydrochloric acid prevents the formation of ferric hydroxide, and
so far reduces the ionization of the water that the hydrolysis of the
ferric sulphate is also prevented, and no precipitation occurs on
heating.]

[Note 3: The nitric acid, after attaining a moderate strength,
oxidizes the Fe^{++} ions to Fe^{+++} ions with the formation of an
intermediate nitroso-compound similar in character to that formed in
the "ring-test" for nitrates. The nitric oxide is driven out by heat,
and the solution then shows by its color the presence of ferric
compounds. A drop of the oxidized solution should be tested on
a watch-glass with potassium ferricyanide, to insure a complete
oxidation. This oxidation of the iron is necessary, since Fe^{++} ions
are not completely precipitated by ammonia.

The ionic changes which are involved in this oxidation are perhaps
most simply expressed by the equation

3Fe^{++} + NO_{3}^{-}+ 4H^{+} --> 3Fe^{+++} + 2H_{2}O + NO,

the H^{+} ions coming from the acid in the solution, in this case
either the nitric or the hydrochloric acid. The full equation on which
this is based may be written thus:

6FeSO_{4} + 2HNO_{3} + 6HCl --> 2Fe_{2}(SO_{4})_{3} + 2FeCl_{3} + 2NO
+ 4H_{2}O,

assuming that only enough nitric acid is added to complete the
oxidation.]

[Note 4: The ferric hydroxide precipitate tends to carry down some
sulphuric acid in the form of basic ferric sulphate. This tendency is
lessened if the solution of the iron is added to an excess of OH^{-}
ions from the ammonium hydroxide, since under these conditions
immediate and complete precipitation of the ferric hydroxide ensues.
A gradual neutralization with ammonia would result in the local
formation of a neutral solution within the liquid, and subsequent
deposition of a basic sulphate as a consequence of a local deficiency
of OH^{-} ions from the NH_{4